My Chapter 4 Notes from my Ch

1. Chapter 4: Atomic Structure
By Kendon Smith – Columbia Central HS – Brooklyn, MI

2. I. Early Models of the Atom
- All matter is composed of tiny particles called
atoms, which are the smallest particles of an
element that retain its properties and identity
during a chemical reaction.
A. Democritus’s Atomic Philosophy
- Democritus was a Greek philosopher from the
4th century B.C. who first suggested the existence
of tiny particles called “atomos”.
- Democritus believed atoms were indivisible and
indestructible.
- Democritus lacked experimental support for his
ideas.

3. B. Dalton’s Atomic Theory
- John Dalton (1766 – 1844) was an English chemist
and schoolteacher.
- Dalton used experiments to transform Democritus’
ideas into scientific theory.

4. Dalton’s Atomic Theory:
1. All elements are composed of tiny indivisible
particles called atoms.
2. Atoms of the same element are identical. Atoms of
one element are different than atoms of another
element.
3. Atoms of different elements can physically mix
together, or they can chemically combine in simple
whole-number ratios to form compounds.
4. Chemical reactions occur when atoms are
separated, joined, or rearranged, however atoms of
one element are never changed into atoms of a
different element.

6. C. Sizing up the Atom
- Atoms are so tiny that a single copper penny contains
2.4 x 1022 atoms.
- A line of 100,000,000 copper atoms would be
1 centimeter long.
- The radii of most atoms are between 5 x 10-11 m
and 2 x 10-10 m.
Calculate the diameter range of most atoms in pm:
Radius = 5 x 10-11 m Radius = 2 x 10-10 m
Diameter = 1 x 10-10 m Diameter = 4 x 10-10 m
÷ 10-12 = 100 pm ÷ 10-12 = 400 pm

7. II. Structure of the Atom
A. Subatomic Particles
- Atoms are now known to be divisible. They can be
broken down into even smaller particles, called
subatomic particles.
- The three subatomic particles are electrons, protons,
and neutrons.

8. 1. Electrons
a. J. J. Thomson, an English physicist,
discovered the electron in 1897.
b. Electrons are negatively charged subatomic
particles.
c. Thomson performed the cathode ray tube
experiment, in which a beam of negatively
charged particles traveled from the negative
electrode, called the cathode, to the positive
electrode, called the anode.

9. High voltage
anode
cathode

10. d. The ray was deflected by magnets and charged
metal plates.
It was repelled by a negative plate and attracted by
a positive plate.
Because Thomson knew that opposites attract,
electrons must be negative.
e. U.S. physicist Robert Millikan carried out
experiments to measure the mass and charge of
the electron.
- An electron carries exactly one unit of negative
charge = -1.
- An electron’s mass is 1/1840 the mass of a proton ≈
basically ZERO mass.

11. 2. Protons and Neutrons
a. The cathode ray tube experiment taught us some
simple concepts about atoms:
1. Atoms have no net charge; they are electrically
neutral.
2. Electric charges are carried by particles of matter.
3. Electric charges always exist in whole numbers –
no fractions of charge.
4. When equal numbers of negatives and positives join,
particles are neutral.

12. b. This meant there must be a positive particle left
behind when atoms lose their negative charged
electrons!
c. In 1886, Eugen Goldstein discovered positive
particles called protons.
d. In 1932, English physicist James Chadwick
discovered neutrons.
- Neutrons carry no charge and have a mass nearly
equal to a proton.
- Neutrons only contribute mass to an atom, making
some atoms heavier.

13. Relative
Particle Symbol Charge
Mass
1/1840 =
Electron e -
-1
zero!
Proton p+ +1 1
Neutron n0 0 1

14. B. The Atomic Nucleus
1. Rutherford’s Gold Foil Experiment
a. In 1911, Ernest Rutherford tested the current atomic
theory by shooting alpha particles at a very thin
sheet of gold foil.
- Alpha particles are helium atoms that have lost
their electrons.
They are made of two protons and two neutrons,
so they have a double positive charge = +2.
b. It was expected that the alpha particles would pass
through the gold foil but experience some

15. Rutherford’s Gold Foil Experiment

16. c. Surprisingly, a majority of the alpha particles passed
through the gold foil as if there was nothing there,
with a few even bouncing back!
d. This led Rutherford to two important conclusions
about atoms:
1. Atoms are mostly empty space!
(Explains lack of deflections.)
2. All the positive charge and mass of the atom must
be located in a tiny, dense region in the central
core of the atom, called the nucleus.
(Explains occasional bounce backs.)
e. In the nuclear atom, protons and neutrons are
located in the nucleus.
The electrons are distributed in the space around
the nucleus.

17. C. The Bohr Model (from Chapter 5)
1. Niels Bohr improved on Rutherford’s model
of the atom and proposed that the electrons
travel around the nucleus in specific
circular paths, or orbits.
- Orbitals, or energy levels, are larger as you
move further away from the nucleus and can
hold more electrons.
Level 1 = 2 e-
Level 2 = 8 e-
Level 3 = 18 e-
Level 4 = 32 e-

18. Models of the Atom:
Thomson’s “Plum Pudding” Model
Electron (–)
Positive Matrix (+)

19. Models of the Atom:

20. Models of the Atom:
Electron (–)
Cloud
Nucleus (+)
(Protons & Neutrons)

21. Models of the Atom:
The Bohr Model of Electron Orbitals
Electron (–)
Energy Levels
or Orbitals
Nucleus
Protons (+) ++
+ +
Neutrons (0)

22. III. Distinguishing Among Atoms
A. Atomic Number
- Elements are different because they contain
different numbers of protons.
a. Atomic Number = the number of
protons in the nucleus of an atom

23. B. Mass Number
a. Mass Number = the total number of protons and
neutrons in the nucleus
b. Only protons and neutrons add mass to an atom –
Electrons are negligible!
* Mass number is NOT THE SAME as Atomic Mass!
c. Mass numbers are always whole numbers!
d. The number of neutrons is the difference between
mass number and atomic number.
# of neutrons = mass number – atomic number

24. Atomic Number (protons)
Atomic Mass
(not mass #)

25. C. Isotopes
- Atoms of the same element can have different
numbers of neutrons, which gives them different
mass numbers!
a. Isotopes are atoms with the same number of
protons, but different number of neutrons.
b. Isotopes are atoms with the same atomic
number, but different mass numbers.

26. IV. Atomic Mass
A. Atomic Mass Units (amu)
- Even the largest atom is incredibly small!
- A proton has an actual mass of 1.67 x 10-24 grams,
so it is difficult to work with numbers these small.
- Atomic mass units are units of relative mass that
were invented to make the numbers easier to work
with and understand.
- Atomic mass units are based on the mass of an atom
of the isotope Carbon-12, which has a mass of
exactly 12 amu’s.

27. IV. Atomic Mass
A. Atomic Mass Units (amu)
- 1 atom Carbon-12 = 12 amu’s
- Therefore, 1 amu = 1/12th the mass of Carbon-12
- Carbon-12 has 6 protons and 6 neutrons =
12 total particles in the nucleus
- Therefore the mass of 1 p+ or n0 = 1 amu
- What does relative mass tell us? It does not tell us
the actual mass of an atom, but instead it tells us
how it’s mass compares to the standard, which is
Carbon-12.

28. Element Relative Mass Meaning
Magnesium 24 amu 1 atom of Mg is 2x heavier
than an atom of C-12
Helium 4 amu He is 3 times lighter than C
Titanium 48 amu Ti is 4 times heavier than C

29. B. Calculating Atomic Mass Values
- In nature, most elements occur as a
mixture of isotopes
- Each isotope has a different mass number,
so the value used to describe the mass of these
mixed samples is a type of average.
- Average masses are weighted according to
percent abundance, which means that those
isotopes that are more abundant have a greater
influence on the average mass.

30. Example: Chlorine Isotopes

31. NEED TO KNOW:
a. How many isotopes exist for an element?
b. What are the mass numbers of each isotope?
c. What is the percent abundance for each isotope?
Calculation Steps:
1. Multiply each mass number by its % abundance.
(% must be re-written as a decimal!)
2. Add up all the results for the total weighted average
atomic mass.

32. Sample Problems:
14. Boron has two isotopes: boron-10 and
boron-11. Which is more abundant, given
that the atomic mass of boron is 10.81 amu?
10.81
BORON-10 BORON-11
10.5
Straight average? = ________
10.81
Weighted average? = ________ (closer to 11!)

33. Sample Problems:
15. There are three isotopes of silicon; they
have mass numbers of 28, 29, and 30. The
atomic mass of silicon is 28.086 amu.
28 29 30
Weighted Average = 28.086 amu

34. Sample Problems:
16. The element copper has naturally occuring
isotopes with mass numbers of 63 and 65.
The relative abundance values are 69.2% for
63 amu, and 30.8% for 65 amu. Calculate the
average atomic mass of copper.
63 amu x 0.692 = 43.596 ADD
THEM
65 amu x 0.308 = 20.02 UP!
63.616 amu